14.1-14.5 - Standard Electrode Potentials

Redox As Electron Transfer

A redox reaction split into two electron moves

When magnesium is placed in copper(II) sulfate solution, brown copper forms and the blue colour fades. The ionic equation is:

The equation is easier to reason about when it is split into two half-equations:

$\ce{Mg(s) -> Mg2+(aq) + 2e-}$

$\ce{Cu2+(aq) + 2e- -> Cu(s)}$

Magnesium loses electrons, so magnesium is oxidised. Copper(II) ions gain electrons, so copper(II) ions are reduced. The electrons lost in one half-equation are the electrons gained in the other; that electron accounting is the core of redox.

In an electrochemical cell, the two half-reactions are separated so the electrons travel through an external wire. Electron flow is from the half-cell where oxidation occurs to the half-cell where reduction occurs. This lesson does not yet calculate a cell emf, but the direction of electron transfer is already the chemical story behind the voltage.

The same language across the periodic table

Oxidation and reduction must work beyond simple Group 1 or Group 2 reactions. The same electron-transfer model applies across s-, p- and d-block elements:

The species being oxidised supplies electrons to another species, so it acts as a reducing agent. The species being reduced accepts electrons from another species, so it acts as an oxidising agent.

Redox As Oxidation-Number Change

Oxidation number is a bookkeeping model

Electron transfer is often visible in half-equations, but many redox equations are not presented as half-equations. Oxidation numbers let you track the same electron movement in a full equation.

The rule is:

• oxidation is an increase in oxidation number
• reduction is a decrease in oxidation number

For simple ions, the oxidation number is the ionic charge. For covalent substances and complex ions, it is a formal bookkeeping value rather than a claim that the atom actually has that full charge.

In an oxyanion, the oxidation numbers add to the ion’s charge. Oxygen is normally $-2$: for $\ce{MnO4-}$, $x+4(-2)=-1$, giving Mn $+7$. For $\ce{Cr2O7^2-}$, $2x+7(-2)=-2$, giving each Cr $+6$. These are formal oxidation numbers, not charges on separate $\ce{Mn^7+}$ or $\ce{Cr^6+}$ ions in water.

Worked example: p-block and d-block together

Consider this reaction:

Iron changes from $+2$ in $\ce{Fe2+}$ to $+3$ in $\ce{Fe3+}$. The oxidation number has increased, so iron(II) is oxidised:

Chlorine changes from $0$ in $\ce{Cl2}$ to $-1$ in $\ce{Cl-}$. The oxidation number has decreased, so chlorine is reduced:

$\ce{Cl2(aq) + 2e- -> 2Cl-(aq)}$

This is a d-block ion being oxidised by a p-block molecule. The labels still come from the same two tests: electron loss/gain and oxidation-number increase/decrease.

Common-error contrast

Do not decide oxidation and reduction by looking only for oxygen or hydrogen. In A level redox, $\ce{Fe2+ ->Fe3+}$ is oxidation even though oxygen is not involved. The key evidence is the electron lost or the oxidation number increasing from $+2$ to $+3$.

What A Standard Electrode Potential Means

A half-cell cannot be measured on its own

A half-cell contains an oxidised form and a reduced form of a redox system. For example:

$\ce{Cu2+(aq) + 2e- <=> Cu(s)}$

Each half-cell has a tendency to accept or release electrons, but a voltmeter cannot measure an absolute tendency from one half-cell alone. A voltage is a potential difference, so it is always a comparison between two electrode systems.

The hydrogen electrode is assigned $E^{\circ} = 0.00\,\mathrm{V}$, so every other $E^{\circ}$ value is measured relative to that reference.

Reduction potentials in the Data Booklet

The Data Booklet table of standard electrode potentials writes half-equations as reductions. For example:

$\ce{H+(aq) + e- <=> 1/2H2(g)}$ — $E^{\circ} = 0.00\,\mathrm{V}$

$\ce{Zn2+(aq) + 2e- <=> Zn(s)}$ — $E^{\circ} = -0.76\,\mathrm{V}$

$\ce{Cu2+(aq) + 2e- <=> Cu(s)}$ — $E^{\circ} = +0.34\,\mathrm{V}$

$\ce{Fe3+(aq) + e- <=>Fe2+(aq)}$ — $E^{\circ} = +0.77\,\mathrm{V}$

The sign is part of the comparison with the hydrogen electrode. A positive value means the right-hand reduction half-equation has a greater tendency to occur as a reduction than the hydrogen half-equation under standard conditions. A negative value means it has a smaller tendency to occur as a reduction than the hydrogen half-equation under standard conditions.

Do not read an $E^{\circ}$ value as a charge, an energy change, or a concentration. It is a potential difference in volts, measured against an agreed reference electrode.